Method of performing in situ calibrated potentiometric pH measurements

ABSTRACT

A device for the precise and accurate potentiometric pH measurements in situ. Embodiments of a potentiometric device according to the invention consist of one or more glass pH-sensitive electrodes connected to a potentiometer. A key feature of the device is that, rather than being calibrated conventionally with buffers, it can be calibrated with an in situ device that measures pH spectrophotometrically. Spectrophotometric pH measurements obtained via sulfonephthalein absorbance measurements are inherently calibrated (do not require buffers). Thus, devices according to the invention allow for continuous potentiometric pH measurements with occasional spectrophotometric calibrations. The spectrophotometric calibration device consists of a spectrophotometer with associated pumps for combining a sulfonephthalein pH indicator with the aqueous medium whose pH is to be measured. The device will record potentiometric pH measurements for an extended period of time until the spectrophotometric device is autonomously activated for another calibration. In this manner precise and accurate pH measurements can be obtained continuously in the environment, and the low energy expenditure of the potentiometric device provides excellent endurance. Also provided is a method and associated devices for spectrophotometrically determining the salinity of an aqueous medium.

CROSS REFERENCE TO RELATED APPLICATIONS

This application is a continuation-in-part of prior U.S. patentapplication Ser. No. 12/110,730, entitled “Sensor for Direct Measurementof Carbonate Ions in Sea Water” filed on Apr. 28, 2008, which claims thebenefit of priority to U.S. Provisional Patent Application 60/914,384,entitled, “Sensor for Direct Measurement of Carbonate Ions in SeaWater”, filed Apr. 27, 2007, the contents of which applications areherein incorporated by reference.

STATEMENT OF GOVERNMENT INTEREST

This invention was made with Government support under Grant Nos.OCE-0551676 and N00014-03-1-0612 awarded by the National ScienceFoundation and Office of Naval Research, respectively. The Governmenthas certain rights in the invention.

FIELD OF INVENTION

This invention relates to sensors for extended deployment in aquaticenvironments. More specifically, this invention relates to devices andassociated methods for precise and accurate potentiometric pHmeasurements in situ over periods of extended deployment with occasionalspectrophotometric calibrations.

BACKGROUND OF THE INVENTION

Solution pH is widely conceptualized as a master variable in theregulation of natural aqueous systems. It is a key feature indescriptive models of carbonate system chemistry, trace metal speciationand bioavailability, oxidation-reduction equilibria and kinetics,biologically induced carbon system transformations, and the aqueousinteractions and transformations of minerals. Rising levels ofatmospheric CO₂ are leading to ocean acidification. The response ofseawater and freshwaters to acidification processes has created a needfor autonomous global monitoring of ocean water and fresh water pH. Theimportance of pH in investigations of terrestrial and oceanicbiogeochemistry has necessitated improvements in not only the quality ofmeasurements (precision and accuracy), but also the spatial and temporalresolution of measurements in the field. Potentiometric devices arerarely used for in situ environmental pH measurements because in situbuffer calibrations are problematic. Spectrophotometric devices havebeen developed for in situ environmental measurements but the enduranceand measurement frequency of spectrophotometric devices is relativelylow due to high power requirements and the limited longevity of lamps.Achieving meaningful spatial and temporal measurements in the fieldmandates the introduction of robust measurement devices capable ofaccurate and precise measurements over extended timeframes.

SUMMARY OF THE INVENTION

The present invention provides a device capable of precise and accuratepotentiometric pH measurements in situ. Also provided are associatedmethods for in situ calibrated pH measurements. Embodiments of an insitu calibrated potentiometric device according to the invention couplepotentiometric pH measurement systems to systems for calibrating thepotentiometric pH measurement systems. The potentiometric system/aspectof the device employs one or more glass pH-sensitive electrodesconnected to a potentiometer (i.e., a device capable of measuringelectrical potentials [voltages]). A key feature of the potentiometricsystem is that, rather than being calibrated conventionally withbuffers, it can be calibrated with an in situ device that measures pHspectrophotometrically. Spectrophotometric pH measurements obtained viasulfonephthalein absorbance measurements are inherently calibrated anddo not require the buffers necessary for the calibration of typicalpotentiometric devices. Thus, devices according to the invention allowfor continuous potentiometric pH measurements with occasionalspectrophotometric calibrations.

The spectrophotometric calibration system/aspect of the device utilizesa spectrophotometer with associated pumps for combining asulfonephthalein pH indicator with the aqueous medium whose pH is to bemeasured. After a sulfonephthalein indicator and the aqueous medium(e.g., seawater) are combined in an optical cell, light is passedthrough the combined mixture, an absorbance spectrum is obtained andsolution pH is calculated from (a) optical absorbance ratios at multiplewavelengths, (b) temperature, and (c) the salt concentration (e.g.,salinity) of the solution. Through contemporaneous potentiometricmeasurements, the potentiometric pH system (glass electrode plus voltmeter) is calibrated without the use of buffers. The device will recordpotentiometric pH measurements for an extended period of time until thespectrophotometric device is autonomously activated for anothercalibration. In this manner precise and accurate pH measurements can beobtained continuously in the environment, and the low energy expenditureof the potentiometric device provides excellent endurance.

In a first aspect the present invention provides a device for in situcalibrated potentiometric pH measurement. The device includes apotentiometric pH measurement module, a spectrophotometric pHmeasurement module and a calibration module. The calibration module isin communication with the potentiometric pH measurement module and thespectrophotometric pH measurement module. The calibration modulereceives pH calibration data from the potentiometric pH measurementmodule and the spectrophotometric pH measurement module and performscalibrations to the potentiometric pH measurement module using thereceived pH calibration data.

In certain embodiments the received pH calibration data utilizes one ormore substantially contemporaneous pH measurements from thepotentiometric pH measurement module and the spectrophotometric pHmeasurement module. Thus, a potentiometric pH measurement is performedat substantially simultaneously to a spectrophotometric pH measurement.By performing the operations coextensively, fluctuations due to ambientconditions and sample identity are minimized.

In further embodiments of the device the spectrophotometric pHmeasurement module can include a first reservoir containing asulfonephthalein pH indicator, one or more pumps in communication withthe first reservoir and an aqueous medium, a second reservoir to receivecombined aqueous medium/pH indicator (combined mixture), and aspectrophotometer to measure the absorbance characteristics of thecombined mixture. The one or more pumps combine a sulfonephthalein pHindicator with the aqueous medium. In an advantageous embodiment of thedevice the calibration module can include systems for the autonomousactivation of the calibration module. The calibration module can beautonomously activated upon a defined time interval or at an eventtrigger indicating the necessity for calibration of the potentiometer.The potentiometric pH measurement module can include one or more glasspH-sensitive electrodes connected to a potentiometer. The deviceaccording to claim 5 wherein the calibration module is autonomouslyactivated upon a defined time interval or at an event trigger indicatingthe necessity for calibration of the potentiometer.

In certain embodiments the device can include a housing containing thepotentiometric pH measurement module, the spectrophotometric pHmeasurement module and the calibration module in communication. Thehousing can have one or more sample intake ports and one or more sampleexhaust ports to allow the intake and exhaust of the samples. Thehousing can also be sealed to allow operation while partially or totallyimmersed in an aqueous medium.

In a second aspect the present invention provides a device forautonomously-activated in situ calibrated potentiometric pH measurement.The device includes a potentiometric pH measurement module, aspectrophotometric pH measurement module, and an autonomously-activatedcalibration module. The autonomously-activated calibration module is incommunication with the potentiometric pH measurement module and thespectrophotometric pH measurement module. The calibration modulereceives pH calibration data from the potentiometric pH measurementmodule and the spectrophotometric pH measurement module and performscalibrations to the potentiometric pH measurement module using thereceived pH calibration data. The autonomously-activated calibrationmodule can be activated upon a defined time interval or at an eventtrigger indicating the necessity for calibration of the potentiometer.

In certain embodiments of the device the spectrophotometric pHmeasurement module can include a first reservoir containing asulfonephthalein pH indicator, one or more pumps in communication withthe first reservoir and an aqueous medium a second reservoir to receivecombined aqueous medium/pH indicator (combined mixture), and aspectrophotometer to measure the absorbance characteristics of thecombined mixture. The one or more pumps combine a sulfonephthalein pHindicator with the aqueous medium.

In a third aspect the present invention provides a method of performingin situ calibrated potentiometric pH measurements for extended periodsof time. The method includes the steps of recording a plurality ofpotentiometric pH measurements using a potentiometric pH measurementdevice, activating a spectrophotometric calibration system,spectrophotmetrically calibrating the potentiometric pH measurementdevice, and resetting the activator of the spectrophotometriccalibration system to allow the potentiometric pH measurement device toreturn to recording potentiometric pH measurements. Thespectrophotometric calibration system can be autonomously activated. Theautonomous activation can occur upon a defined time interval or at anevent trigger indicating the necessity for calibration.

In certain embodiments the step of spectrophotmetrically calibrating thepotentiometric pH measurement device includes the steps of sampling anaqueous medium, accessing data representative of the temperature andsalt concentration of the sampled aqueous medium, combining asulfonephthalein indicator and a first portion of the sampled aqueousmedium, delivering the combined medium to an optical cell, obtaining anabsorbance spectrum from the combined medium, calculating solution pH ofthe sampled aqueous medium utilizing the optical absorbance ratios atmultiple wavelengths, the temperature, and the salt concentration,obtaining a potentiometric pH measurement of a second portion of theaqueous medium, calculating the deviation between the potentiometric pHmeasurement and the spectrophotometric pH measurement, and adjusting thepotentiometric pH measurement device using the calculated deviation. Thespectrophotometric pH measurement is considered accurate, while anydeviation between the potentiometric pH measurement and thespectrophotometric pH measurement is due to discrepencies in thepotentiometric pH measurement due to factors such as the inherent driftobserved in this type of unit over time. The adjusting step calibratesthe potentiometric pH measurement device.

In certain embodiments the step of obtaining a potentiometric pHmeasurement of a second portion of the aqueous medium is performedsubstantially contemporaneously with the step of obtaining an absorbancespectrum from the combined medium.

In a fourth aspect the present invention provides a method ofspectrophotmetrically calibrating a potentiometric pH measurementdevice. The method includes the steps of obtaining one or more samplesof an aqueous medium, accessing data representative of the temperatureand salt concentration of the sampled aqueous medium, combining asulfonephthalein indicator and a first portion of the sampled aqueousmedium, delivering the combined medium to an optical cell, obtaining anabsorbance spectrum from the combined medium and calculating solution pHof the sampled aqueous medium utilizing the optical absorbance ratios atmultiple wavelengths, the temperature, and the salt concentration,obtaining a potentiometric pH measurement of a second portion of theaqueous medium, calculating the deviation between the potentiometric pHmeasurement and the spectrophotometric pH measurement, and adjusting thepotentiometric pH measurement device using the calculated deviation. Theadjusting step calibrates the potentiometric pH measurement device.

In certain embodiments the step of obtaining a potentiometric pHmeasurement of a second portion of the aqueous medium is performedsubstantially contemporaneously with the step of obtaining an absorbancespectrum from the combined medium. In additional embodiments the methodof spectrophotmetrically calibrating the potentiometric pH measurementcan be autonomously activated. Autonomous activation can occurs upon adefined time interval or at an event trigger indicating the necessityfor calibration or under other conditions.

In a fifth aspect the present invention provides a method of measuringthe salinity of an aqueous solution, the method includes the steps ofobtaining a sample of an aqueous medium, adding a metal ion species theaqueous medium to produce a sample solution, spectrophotometricallymeasuring the absorbance of light passing through the sample solution ata plurality of wavelengths, wherein the ultraviolet absorbance of lightis a function of the complexation in the sample solution of the addedmetal ion species with carbonate ions of the sample and computing thesalinity of the aqueous medium based upon the absorbance ratio of thesample solution at the plurality of wavelengths.

In certain embodiments the spectrophotometric absorbance measurementsare obtained in the ultraviolet range.

In certain embodiments the metal ion can include Pb^(II) and Cu^(II),yttrium, a lanthanide metal and an actinide metal. In embodiments wherethe metal ion species is Pb^(II), one of the pluralityspectrophotometric absorbance measurements used to calculate theabsorbance ratio can be measured at about λ=234 nm In further aspects ofthis embodiment a second of the plurality of spectrophotometricabsorbance measurements used to calculate the absorbance ratio ismeasured at about λ=240 nm. to about λ=260 nm In still further aspects,the metal ion species is Pb^(II) and the spetrophotometric absorbancemeasurements are measured at about λ=234 nm. and about λ=250 nm.

BRIEF DESCRIPTION OF THE DRAWINGS

For a fuller understanding of the invention, reference should be made tothe following detailed description, taken in connection with theaccompanying drawings, in which:

FIG. 1 is a graph showing Pb(II) UV absorbance spectra in seawater atS=35.87 and 25° C. as a function of pH.

FIG. 2 is a graph showing Pb(II) UV absorbance spectra of acidifiedseawater (25° C., pH=3.73) at four selected salinities.

FIG. 3 is a graph showing salinity dependence of PbCO₃ ⁰ formationconstant at 25° C.

FIG. 4 a is a graph showing salinity dependence of ₂₅₀ε_(Pb) and₂₅₀ε_(PbCO3) at 25° C.

FIG. 4 b is a graph showing salinity dependence of ₂₃₄ε_(Pb) and₂₃₄ε_(PbCO3) at 25° C.

FIG. 5 is a graph showing Pb(II) absorbance ratios of in acidifiedseawater at 250 and 234 nm: ₂₅₀A/₂₃₄A (₂₅₀ε_(Pb)/₂₃₄ε_(Pb)) as afunction of salinity.

FIG. 6 is a series of three graphs showing salinity dependence of e₁ (6a), e₂ (6 b), and e₃ (6 c) at 25° C.

FIG. 7 a is a graph showing the best fit log {(_(CO3)β₁)/(e₂)} resultsusing Eq. (20)

FIG. 7 b is a graph showing the best fit e₁ results using Eq. (20).

FIG. 8 is a graph showing residuals ([CO₃ ²⁻]_(observed)−[CO₃²]_(predicted)) plotted as a function of ([CO₃ ²⁻]_(observed) for theleast squares analyses using Eq. (20).

DETAILED DESCRIPTION OF THE PREFERRED EMBODIMENT

Rising levels of atmospheric CO₂ are leading to ocean acidification. Theresponse of seawater and freshwaters to acidification processes hascreated a need for autonomous global monitoring of ocean water and freshwater pH. Potentiometric devices are rarely used for in situenvironmental pH measurements because in situ buffer calibrations areproblematic. Spectrophotometric devices have been developed for in situenvironmental measurements but the endurance and measurement frequencyof spectrophotometric devices is relatively low due to high powerrequirements and the limited longevity of lamps. Spectrophotometric andpotentiometric devices can be combined to reduce power, enhancelongevity and still provide high quality calibrated measurements.

Both potentiometric and spectrophotometric procedures can be used for pHmeasurements. Potentiometric pH measurements can be performed withrelatively simple equipment and procedures required. This simplicitymakes potentiometry a good choice for field measurements as long asthere are not stringent requirements for accuracy and precision. Underideal conditions, potentiometric measurements that utilize glasshydrogen ion selective electrodes can provide measurement precisions onthe order of 0.003 pH units. However, measurement accuracy can beproblematic. Potentiometric measurements generally require regularbuffer calibrations, and special care must be taken to address artifactsassociated with both residual liquid junction potentials and variationsin asymmetry potentials. Seiter and DeGrandpre evaluated performance ofsix electrodes under identical operational conditions. They observedthat individual electrodes generally have distinctive drift patterns,with drift rates up to 0.02 pH units per day (Seiter, J. C.; DeGrandpre,M. D. Talanta 2001, 54, 99). Electrode drift necessitates frequentcalibrations, making autonomous operation somewhat problematic comparedto spectrophotometric pH determinations. One important advantage ofpotentiometric devices is their low power consumption. This aspect wouldseem to make them well-suited for deployment over extended periods oftime. Unfortunately, their use for such deployment is greatly limited bytheir inherent drift, making readings increasingly inaccurate andnecessitating calibration for meaningful use.

Although potentiometric pH measurements are versatile and satisfactoryfor many applications, spectrophotometric pH measurement procedures haveat least two important advantages that make them particularly desirable.Since spectrophotometric pH measurements can be determined viaabsorbance ratios, and the calibration of pH indicators is a laboratoryexercise that establishes how each indicator's molecular properties varywith temperature, pressure and ionic strength, spectrophotometric pHmeasurements are inherently calibrated and can be termed “calibrationfree”. Subsequent to careful laboratory calibration, spectrophotometricpH measurements do not require the use of buffers. Secondly, thousandsof at-sea observations have demonstrated that the imprecision ofshipboard spectrophotometric pH measurements is on the order of 0.0003to 0.0004 pH units, approximately an order of magnitude better thanpotentiometric results. These advantageous attributes ofspectrophotometric pH measurements have made spectrophotometricprocedures valuable for not only observations of pH, but also formeasurements of CO₂ fugacity and total dissolved inorganic carbon.However, the power consumption of spectrophotometers is considerablygreater than potentiometers. This can be a limitation for theirlong-term deployment. The limited longevity of lamps inspectrophotometers is an additional impediment to their long-termdeployment.

Spectrophotometric pH measurements have been increasingly utilized formeasurements of pH in natural waters. Bellerby et al. developed a flowinjection procedure for spectrophotometric measurement of seawater pHwith a reported precision of 0.005 pH units and a sample frequency of 25hr⁻¹ (Bellerby R. G. J.; Turner, D. R.; Millward, G. E.; Worsfold P. J.Analytica Chimica Acta 1995, 309, 259.). Tapp et al. described the useof a shipboard system for spectrophotometric measurements of surfacewater pH with a reported precision on the order of 0.001 pH units and a1-Hz measurement frequency (Tapp, M.; Hunter, K.; Currie, K.;Mackaskill, B. Mar. Chem. 2000, 72, 193.). Relative to discretemeasurements however, observed discrepancies were as large as 0.02 pHunits. Martz et al. described the construction of a submersible pHsensor with a 0.003 unit measurement precision and a measurementfrequency of 6 hr⁻¹ (Martz, T. R.; Carr, J. J.; French, C. R.;DeGrandpre, M. D. Anal. Chem. 2003, 75, 1844.).

Submersible potentiometric sensors according to the present inventionprovide continuous unattended measurements of seawater pH for periods ofsix to twelve months with a measurement frequency up to 60 Hz. Nocalibrations are required on the part of the user. Both measurements andcalibrations are performed in-situ. Measurement accuracy is on the orderof 0.001 pH units. High measurement frequency, excellent endurancecharacteristics, and in situ calibration make this sensor system wellsuited to descriptions of acid-base equilibrium and kinetic phenomena ondiurnal, seasonal and annual scales.

The present invention provides a device capable of precise and accuratepotentiometric pH measurements in situ. Embodiments of a potentiometricdevice according to the invention consist of one or more glasspH-sensitive electrodes connected to a potentiometer (e.g., a devicecapable of measuring electrical potentials [voltages]). A key feature ofthe device is that, rather than being calibrated conventionally withbuffers, it can be calibrated with an in situ device that measures pHspectrophotometrically. Spectrophotometric pH measurements obtained viasulfonephthalein absorbance measurements are inherently calibrated (donot require buffers). Thus, devices according to the invention allow forcontinuous potentiometric pH measurements with occasionalspectrophotometric calibrations. The spectrophotometric calibrationdevice consists of a spectrophotometer with associated pumps forcombining a sulfonephthalein pH indicator with the aqueous medium whosepH is to be measured. After a sulfonephthalein indicator and the aqueousmedium (e.g., seawater) are combined in an optical cell, light is passedthrough the combined mixture, an absorbance spectrum is obtained andsolution pH is calculated from (a) optical absorbance ratios at multiplewavelengths, (b) temperature, and (c) the salt concentration (e.g.,salinity) of the solution. Through contemporaneous potentiometricmeasurements the potentiometric pH system (glass electrode plus voltmeter) is calibrated without the use of buffers. The device will recordpotentiometric pH measurements for an extended period of time until thespectrophotometric device is autonomously activated for anothercalibration. In this manner precise and accurate pH measurements can beobtained continuously in the environment, and the low energy expenditureof the potentiometric device provides excellent endurance.

It is contemplated that calibrations can be triggered by a variety ofevents or circumstances. For example, calibrations can be triggeredbased upon time intervals or other user-defined events. Alternatively,or in conjunction with such tachniques, calibrations can be triggeredbased upon observations of potentiometric pH changes. For example, asudden pH cahnge outside of the typical deviations of pH recordingscould trigger a calibration to ensure that the large recorded deviaitonwas real.

A number of other solutions are available to avoid problems associatedwith potiometric drift in the long-term deployment of these devices.Pontiometric pH measurement devices consist of (a) a glass pH sensitivemembrane and (b) a reference electrode. One of the main sources problemsin a potentiometric pH measurement is clogging of what is called the‘liquid junction’ of the reference cell. This is an event that couldinitiate drift. If the potentiometric system is initially calibrated,some slow drift may be acceptable because subsequent deviations betweenspectrophotometric and potentiometric measurements might be small. Ifdeviations suddenly become large, the device could be told to select adifferent reference electrode (or a different glass electrode).Alternatively, or in conjunctions with the above procedures, the systemcould be directed to flush itself with a microbial inhibitor or toactivate an ultrasonic device for cleaning.

Procedures for direct measurements of carbonate ion concentrations andsaturation states in seawater have also been developed. Measurements areobtained via ultraviolet spectroscopic observations of Pb(II) spectra asthe relative concentrations of PbCO₃ ⁰ and an ensemble of lead chloridecomplexes vary in response to dissolved CO₃ ²⁻. Measurement precision isenhanced by parameterization in terms of absorbance ratios. The PbCO₃ ⁰stability constant, and Pb(II) molar absorbance ratios in seawater, weredetermined at 25° C. over a range of salinity between 36 and 20. Theprocedures described herein are well suited to measurements throughoutthe normal range of carbonate ion concentrations in the oceans. Rapidequilibration rates for Pb(II) carbonate complexation make theprocedures described in this work well suited to rapid direct analysisin situ.

The health of coral reefs and calcareous plankton is also stronglyinfluenced by the carbonate saturation state of seawater. Calculationsof carbonate saturation states currently require measurements of two CO₂system parameters, such as pH and total dissolved carbon, plusthermodynamic calculations that relate carbonate ion concentrations todirectly measured parameters.

Investigations of the marine CO₂ system are commonly conducted throughmeasurements of four primary variables: total dissolved inorganic carbon(C_(T)), total solution alkalinity (A_(T)), CO₂ fugacity (_(f)CO₂), andsolution pH. Thermodynamic models link these four primary variables,whereby measurements of any two variables can be used to calculate thetwo remaining parameters (DOE, 1994). These models also allowcalculations of the concentrations of the individual forms of inorganiccarbon in seawater: the dissolved concentrations of CO₂ and H₂CO₂, andthe free plus ion paired concentrations of bicarbonate, HCO₃ ⁻, andcarbonate, CO₃ ²⁻. Two of the directly measured and derived CO₂ systemvariables can be highlighted because of their special significance inevaluations of global carbon fluxes and the biogeochemistry of marinecarbonates in general. CO₂ fugacity measurements are essential todescriptions of CO₂ exchange at the air sea interface (DOE, 1994;McGillis and Wanninkhof, 2006; Millero, 2007), and carbonate ionconcentrations are essential to evaluations of (a) the mineralizationrates of marine calcifiers (Langdon and Atkinson, 2005) and (b) thedissolution rates of calcite and aragonite (CaCO_(3(s)) polymorphs) bothon the seafloor and in the water column (Morse, 1978; Keir, 1980; Ackeret al., 1987).

Rising atmospheric carbon dioxide concentrations over the past twocenturies have led to increasing CO₂ uptake by the oceans (RoyalSociety, 2005). This process, which is decreasing the pH of the upperocean, is reducing oceanic carbonate ion concentrations and thus thelevel of saturation of calcium carbonate (Broecker et al, 1979; Feely etal., 2004; On et al., 2005). If the trend continues, it will have aseriously negative impact on key marine organisms such as corals andsome plankton (Kleypas et al, 2006). In view of the importance ofcarbonate ion concentrations ([CO₃ ²⁻]_(T)) to the oceans' rapidlyevolving carbonate system, it is then highly desirable to move [CO₃²⁻]_(T) from the rank of derived CO₂ system variables to the list ofprimary measured variables.

A variety of metals in seawater, including lead (Byrne, 1981), copper(Byrne and Miller, 1985), the lanthanides (Cantrell and Byrne, 1987),and various actinides (Byrne, 2002) have inorganic speciation schemesthat are strongly dominated by carbonate complexation. Among thesemetals, the speciation of Pb(II) and Cu(II) has been examined directlyby ultraviolet absorbance spectroscopy in natural seawater (Byrne, 1981;Byrne and Miller, 1985). Since the ultraviolet absorbancecharacteristics of these metals are strongly influenced by dissolvedcarbonate, it follows that observations of Pb(II) and Cu(II) absorbancespectra can be used to directly determine seawater carbonate ionconcentrations. As a means of achieving high precisions in suchdeterminations, we have developed procedures that involve measurementsof absorbance ratios rather than absolute absorbance. Our techniques areclosely analogous to those that were developed previously for seawaterpH measurements with precisions on the order of 0.0004 pH units (RobertBaldo et al, 1985; Byrne, 1987; Byrne and Breland, 1989; Clayton andByrne, 1993). Whereas spectrophotometric observations ofsulfonephthalein acid/base equilibria are utilized for seawater pHmeasurements, spectrophotometric observations of metal ion complexationcan be used to quantify anion concentrations in seawater. Lead isespecially well suited to such measurements because (a) PbCO₃ ⁰ and avariety of Pb(II) chloride complexes have dissimilar absorbance spectrain the ultraviolet, and (b) species other than PbCO₃ ⁰ and chloridecomplexes appear to be insignificant over a wide range of salinities innatural seawater. In this work, Pb(II) formation constants and molarabsorbance ratios required for direct determinations of carbonate ionconcentrations in seawater are characterized at 25° C. as a function ofsalinity. In addition to development of procedures for measurement ofcarbonate ions and carbonate saturation state, we also show thatmeasurements of Pb(II) absorbance ratios in acidified seawater can beused to determine seawater salinity with a precision on the order of±0.1 salinity units. The procedures described in this work are suitablefor rapid autonomous in-situ monitoring of carbonate ion concentrationin seawater.

Theoretical Principles

The PbCO₃ ⁰ formation reaction in seawater,Pb²⁺+CO₃ ²⁻

PbCO₃ ⁰,  (1)can be quantitatively described with an equilibrium constant of thefollowing form:

$\begin{matrix}{{{}_{{CO}\; 3}^{}{}_{}^{}} = \frac{\left\lbrack {PbCO}_{3}^{0} \right\rbrack_{T}}{{\left\lbrack {Pb}_{T} \right\rbrack\left\lbrack {CO}_{3}^{2 -} \right\rbrack}_{T}}} & (2)\end{matrix}$where [Pb_(T)] represents the total concentration of Pb(II) speciesother than PbCO₃ ⁰ in seawater, principally Pb²⁺, PbCl⁺, PbCl₂ ⁰ andPbCl₃ ⁻, and minor amounts of PbSO₄ ⁰; [CO₃ ²⁻]_(T) is the sumconcentration of free and ion paired carbonate (CO₃ ²⁻, NaCO₃ ⁻, MgCO₃ ⁰and CaCO₃ ⁰; and [PbCO₃ ⁰]_(T) represents the sum concentration of PbCO₃⁰ and potentially significant mixed ligand complexes such as PbCO₃Cl⁻.The absorbance of Pb(II) in seawater can be described using thefollowing equation (Byrne, 1981; Soli et al, 2008):

$\begin{matrix}{\frac{\,_{\lambda}A}{l \cdot \lbrack{Pb}\rbrack_{T}} = \frac{{{}_{}^{}{}_{}^{}} + {{{}_{}^{}{}_{{PbCO}\; 3\mspace{11mu}{CO}\; 3}^{}}{\beta_{1}\left\lbrack {CO}_{3}^{2 -} \right\rbrack}_{T}}}{1 + {{{}_{{CO}\; 3}^{}{}_{}^{}}\left\lbrack {CO}_{3}^{2 -} \right\rbrack}_{T}}} & (3)\end{matrix}$where _(λ)A is the absorbance of Pb(II) at wavelength λ, l is thepathlength, [Pb]_(T) is the total lead concentration, _(λ)ε_(PbCO3) isthe molar absorbance of)(PbCO₃ ⁰)_(T) at wavelength λ, _(λ)ε_(Pb) is themolar absorbance of (Pb_(T)) at wavelength λ and _(CO3)β₁ is theformation constant of PbCO₃ ⁰ as defined in Eq. (2). Eq. (3) can be usedto describe the dependence of Pb(II) absorbance (_(λ)A) on [CO₃ ²⁻]_(T)and determine an internally consistent set of values for_(λ)ε_(Pb,λ)ε_(PbCO3) and _(CO3)β₁.

Use of Eq. (3) at wavelengths and allows carbonate ion concentrations,[CO₃ ²⁻]_(T), to be directly calculated from observations of absorbanceratios:

$\begin{matrix}{R = {\frac{\,_{\lambda 2}A}{\,_{\lambda 1}A} = \frac{{{}_{}^{}{}_{}^{}} + {{{}_{}^{}{}_{{PbCO}\; 3\mspace{11mu}{CO}\; 3}^{}}{\beta_{1}\left\lbrack {CO}_{3}^{2 -} \right\rbrack}_{T}}}{{{}_{}^{}{}_{}^{}} + {{{}_{}^{}{}_{{PbCO}\; 3\mspace{11mu}{CO}\; 3}^{}}{\beta_{1}\left\lbrack {CO}_{3}^{2 -} \right\rbrack}_{T}}}}} & (4)\end{matrix}$Rearrangement of Eq. (4) results in the following equation:

$\begin{matrix}{{- {\log\left\lbrack {CO}_{3}^{2 -} \right\rbrack}_{T}} = {{\log_{{CO}\; 3}\beta_{1}} + {\log\left( \frac{R - e_{1}}{e_{2} - {R \cdot e_{3}}} \right)}}} & (5)\end{matrix}$where e₁, e₂, and e₃ are Pb(II) molar absorbance ratios:e ₁=_(λ2)ε_(PbCO3)/_(λ1)ε_(PbCO3) ,e ₂=_(λ2)ε_(Pb)/_(λ1)ε_(PbCO3) ,e₃=_(λ1)ε_(PbCO3)  (6)

The form of Eq. (5) is identical to that which has been used for highlyprecise measurements of seawater pH from observations ofsulfonephthalein absorbance in seawater (Robert Baldo et al, 1985;Byrne, 1987; Byrne and Breland, 1989; Clayton and Byrne, 1993).

At sufficiently low pH (i.e., where [CO₃ ²⁻]_(T)˜0), Eq. (4) can bewritten as:R= _(λ2) A/ _(λ1) A= _(λ2)ε_(Pb)  (7)It has been shown (Byrne et al., 1981) that characterizations of themolar absorptivities of individual species of Pb(II) can be used todirectly determine the relative concentrations of Pb²±, PbCl⁺, PbCl₂ ⁰and PbCl₃ ⁻ in both synthetic solutions and seawater. Since the relativeconcentrations of these species are directly dependent on the chlorideconcentrations in synthetic solutions and seawater, it follows thatPb(II) absorbance ratios at low pH are directly dependent on salinity.In addition to developing a direct means of determining carbonate ionconcentrations via Eq. (5), we show in this work that observations ofPb(II) absorbance ratios at low pH allow calculations of seawatersalinity with a precision somewhat better than ±0.2%.Methods

Eq. (5) can be used to determine carbonate ion concentrations via directmeasurements of Pb(II) absorbance ratios, and characterizations of_(CO3)β₁, e₁, e₂, and e₃. Observations of Pb(II) absorbance spectra atsalinities typical of open ocean seawater (S=35.87) reveal isosbesticpoints near 234 nm (FIG. 1). On this basis, one of the two wavelengthschosen for absorbance observations was λ₁=234 nm Although use of shorterwavelengths is desirable as a means of increasing sensitivity toformation of PbCO₃ ⁰, small absorbance contributions from carbonate ionsat shorter wavelengths make interpretations of absorbances at λ<234 nmless direct. In view of the substantial absorbance variations between240 and 260 nm (FIG. 1), the second of the two wavelengths chosen forabsorbance ratio observations was λ₂=250 nm.

Characterizations of log _(CO3)β₁ in this work were obtained using Eq.(3) and measurements of ₂₅₀A and [CO₃ ²⁻]_(T) in seawater samples atconstant salinity and constant temperature. Along with characterizationsof log _(CO3)β₁, these measurements also produced pairedcharacterizations of ₂₅₀ε_(Pb) and ₂₅₀ε_(PbCO3). Pairedcharacterizations of ₂₃₄ε_(Pb) and ₂₃₄ε_(PbCO3) were obtained frommeasurements of ₂₃₄A and [CO₃ ²⁻]_(T) using Eq. (3), and the log_(CO3)β₁ values determined at each salinity, as described above. Pairedcharacterizations of ₂₃₄ε_(Pb) and ₂₅₀ε_(Pb) were obtained from _(λ)Aobservations at low pH (FIG. 2). The molar absorbance ratios e₁, e₂ ande₃ in Eq. (5) were then determined from these paired molar absorbancecharacterizations as follows:e₁=(₂₅₀ε_(PbCO3)/₂₅₀ε_(Pb))×(₂₃₄ε_(Pb)/₂₃₄ε_(PbCO3))×(₂₅₀ε_(Pb)/₂₃₄ε_(Pb))  (8)e ₂=(₂₃₄ε_(Pb)/₂₃₄ε_(PbCO3))×(₂₅₀ε_(Pb)/₂₃₄ε_(Pb))  (9)e ₃=(₂₃₄ε_(Pb)/₂₃₄ε_(PbCO3))  (10)

All chemicals used were analytical reagent grade. PbC₁₂ and NaHCO₃ werefrom Sigma-Aldrich. HCl (1.000 M) was from J. T. Baker. The seawaterused in this study was surface water from the Gulf of Mexico. Seawatersalinity was measured with an SBE 49 CTD (Seabird). Seawater samples atvarious salinities were prepared by dilution with Milli-Q water.Absorbance measurements were obtained using quartz optical cells (10 cmpathlength) in an HP 8453 spectrophotometer. The slitwidth of thisspectrophotometer is 1 nm. Use of a spectrophotometer with asubstantially larger slitwidth can alter the wavelength-dependentabsorbance characteristics of the parameters given in equations(8)-(10). The temperature of the samples in the optical cells wascontrolled (25±0.05)° C. with a Neslab refrigerating circulator and awater-jacketed spectrophotometric cell holder.

Seawater alkalinity was determined using a spectrophotometric procedure(Yao and Byrne, 1998) that is precise to better than 1 mmol/kg. Seawater(140.0 g) was added gravimetrically to an open top optical cell (10 cmpathlength) which, in turn, was positioned in the thermostated cellholder. Sample pH was measured using an Orion Ross-type pH electrode(No. 800500) connected to an Orion pH meter (Model 720A) in the absolutemillivolt mode. Nerstian behavior of the pH electrode was confirmed viatitrations of 0.7 molal NaCl solutions with concentrated HCl. Theelectrode was calibrated on the total hydrogen ion concentration scalethrough measurements in natural seawater whose pH was determined bysimultaneous spectrophotometric observations of thymol blue absorbanceratios (Zhang and Byrne, 1996).

Through addition of NaHCO₃, the alkalinity of each seawater sample wasincreased to values approximately double those of natural seawater(final alkalinity ˜4.0 millimolal). Seawater samples had CO₂ fugacitiesgenerally in excess of 500 μatm and pH≦8.0. After each seawater samplewas thermally equilibrated, a reference spectrum was taken and 1.05 mlof a 0.001 mol/kg PbCl₂ stock solution was added to the sample (final[Pb(II)]_(T)˜7.5 μmol kg⁻¹). An absorbance spectrum was then taken alongwith a potentiometric measurement of pH. The sample was subsequentlytitrated with standard HCl using a Gilmont micrometer syringe. HCladditions were quantified gravimetrically. Pb(II) absorbance, alkalinityand pH were recorded for each titration point. Sample alkalinity wascalculated by accounting for cumulative HCl additions to the initialseawater sample. Calculations of [CO₃ ²⁻]_(T) from alkalinity and pHutilized the total H⁺ scale dissociation constants of Dickson andMillero (1987) that were derived from the data of Mehrbach et al.(1973). All such calculations were performed using the CO₂ systemprogram of Pierrot et al. (2006). Based on 95% confidence intervals fortotal alkalinities on the order of 0.1%, and 95% confidence intervalsfor open-cell pH measurements on the order of 0.01 units, correspondingrelative errors in calculated carbonate ion concentrations areapproximately 2.3% (e.g., ±5.8 μmol/kg when [CO₃ ²⁻]_(T)=250 μmol/kg).Non-linear least squares parameter estimates of ₈₀ ε_(Pb,λ)ε_(PbCO3) and_(CO3)β₁ were obtained using Eq. (3) and paired values of _(λ)A and [CO₃²⁻]_(T). Calculations of [CO₃ ²⁻]_(T) that accounted for minorcontributions of PbCO₃ ⁰ to carbonate alkalinity did not causesignificant changes in derived values of _(λ)ε_(Pb,λ)ε_(PbCO3) and_(CO3)β₁. Absorbance contributions of CO₃ ²⁻ at short wavelengths wereexamined by performing titration experiments without addition of Pb(II)to samples. Observations of well defined isosbestic points near λ=234 nmdemonstrate that these corrections are very small at the wavelengthsutilized in this work. The dependencies of e₁, e₂ and e₃ on salinitywere described via quadratic functions.

Pb(II) absorbance measurements in acidified seawater (pH ˜3.7,[Pb(II)]_(T) ˜7.5 μmol kg⁻¹) were used to determine _(λ)ε values atλ=234 and 250 nm. Absorbances, in this case, were measured against areference solution of acidified seawater that contained no lead. Theabsorbance ratios obtained in these experiments (₂₅₀ε_(Pb)/₂₃₄ε_(Pb))were used in determinations of e₁ and e₂, as described above, and werealso used in a least squares quadratic regression that allows salinity(S) to be calculated from ₂₃₄A/₂₅₀A observations at low pH.

Results and Discussion

Salinity Dependencies of Pb(II) Molar Absorptivities and the PbCO₃ ⁰Formation Constant

Estimates for _(CO3)β_(1,250)ε_(Pb) and ₂₅₀ε_(PbCO3) obtained using Eq.(3) are given in Table 1 and are shown graphically in FIGS. 3 and 4 a.Over a salinity range between S=20 and S=36, the dependence of the PbCO₃⁰ formation constant on S at 25° C. (FIG. 3) can be described as:log _(CO3)β₁=6.574−0.1235S+1.514×10⁻³ S ²  (11)with a ±0.023 standard error of estimation. The ₂₅₀ε_(Pb), and₂₅₀ε_(PbCO3) values determined in this analysis exhibited a lineardependence on salinity:₂₅₀ε_(Pb)=8.443×10⁵+7.258×10⁴ S  (12)₂₅₀ε_(PbCO3)=4.563×10⁵+2.252×10⁴ S  (13)

The ₂₃₄ε_(Pb) and ₂₃₄ε_(PbCO3) values determined using the log _(CO3)β₁results in Table 1 and absorbance observations at 234 nm are given inTable 2. The FIG. 4 b graphical depiction of these results shows thatthe dependence of ₂₃₄ε_(Pb) on salinity is linear while satisfactorydescriptions of ₂₃₄E_(PbCO3) require a quadratic term:₂₃₄ε_(Pb)=3.837×10⁶+1.749×10³ S  (14)₂₃₄ε_(PbCO3)=3.055×10⁵+2.028×10⁵ S−2.548×10³ S ²  (15)

TABLE 1 PbCO₃ ⁰ formation constant (_(CO3)β₁), ₂₅₀ε_(Pb), and₂₅₀ε_(PbCO3) as function of salinity at 25° C. ₂₅₀ε_(Pb) ₂₅₀ε_(PbCO3)Salinity log _(CO3)β₁ (×10⁻⁶ cm² mol⁻¹) (×10⁻⁶ cm² mol⁻¹) 35.87 4.070 ±0.030 3.445 ± 0.028 1.200 ± 0.028 34.50 4.140 ± 0.028 3.339 ± 0.0351.297 ± 0.024 32.50 4.151 ± 0.026 3.176 ± 0.028 1.193 ± 0.025 30.004.249 ± 0.021 3.012 ± 0.021 1.167 ± 0.019 27.50 4.290 ± 0.024 2.900 ±0.028 1.027 ± 0.019 25.00 4.443 ± 0.028 2.703 ± 0.032 1.024 ± 0.01722.50 4.559 ± 0.018 2.455 ± 0.020 0.976 ± 0.008 20.00 4.711 ± 0.0222.263 ± 0.021 0.899 ± 0.011

The salinity dependence for observations of ₂₅₀ε_(Pb)/₂₃₄ε_(Pb) inacidified seawater (Table 3 and FIG. 5) is well described by thefollowing expression:₂₅₀ε_(Pb)/₂₃₄ε_(Pb)=0.1931+2.062×10⁻² S−3.852×10⁻⁵ S ²  (16)

Using the results that are summarized in Eqs. (12) through (16), Eqs.(8) through (10) can be used to calculate e₁, e₂ and e₃ at eachsalinity. The coefficients obtained in this manner are given in Table 4and are depicted graphically in FIGS. (6 a), (6 b) and (6 c).

The salinity dependencies of e₁, e₂ and e₃ are then given as follows:e ₁=0.3447−6.662×10⁻³ S+1.463×10⁻⁴ S ²  (17)e ₂=0.7749−1.122×10⁻² S+3.331×10⁴ S ²  (18)e ₃=2.114−6.600×10⁻² S+9.036×10⁻⁴ S ²  (19)

TABLE 2 ₂₃₄ε_(Pb) and ₂₃₄ε_(PbCO3), as a function of salinity at 25 ° C.Salinity ₂₃₄ε_(Pb) (×10⁻⁶ cm² mol⁻¹) ₂₃₄ε_(PbCO3) (×10⁻⁶ cm² mol⁻¹)35.87 3.904 ± 0.188 4.280 ± 0.029 34.50 3.881 ± 0.155 4.253 ± 0.03532.50 3.859 ± 0.171 4.255 ± 0.008 30.00 3.872 ± 0.129 4.135 ± 0.00727.50 3.959 ± 0.153 3.899 ± 0.012 25.00 3.939 ± 0.148 3.759 ± 0.01622.50 3.848 ± 0.076 3.595 ± 0.029 20.00 3.832 ± 0.104 3.348 ± 0.047

TABLE 3 Absorbance of Pb(II) in acidified seawater (pH = 3.73):₂₅₀A/₂₃₄A = ₂₅₀ε_(Pb)/₂₃₄ε_(Pb) as function of salinity at 25° C.Salinity ₂₅₀A/₂₃₄A (₂₅₀ε_(Pb)/₂₃₄ε_(Pb)) 20.00 0.5906 21.25 0.6135 22.500.6380 23.73 0.6593 25.00 0.6863 26.25 0.7069 27.50 0.7327 28.75 0.752630.00 0.7778 31.24 0.7987 32.50 0.8231 33.75 0.8448 34.50 0.8603 35.000.8688 35.61 0.8788 35.87 0.8830 35.87 0.8818Determinations of CO₃ ²⁻ Concentrations in Seawater

Eq. (5), Eq. (11), and Eqs. (17) through (19) permit direct measurementsof [CO₃ ²⁻]_(T) from measurements of Pb(II) absorbance ratios inseawater at 25° C. Eq. (5) can, however, also be written in analternative form, with a smaller number of parameters:−log [CO₃ ²⁻]_(T)=log {(_(CO3)β₁)/(e₂)}+log {(R−e₁)/(1−Re₃/e₂)}  (20)

This equation is advantageous for calculations of carbonate ionconcentrations because (a) it reduces the number of parameterizationsrequired for measurements—Using Eq. (20), (_(CO3)β₁)/(e₂) is determinedas a single parameter, and (e₃/e₂) is determined as a single parameter;(b) the parameter (e₃/e₂) can be precisely determined from directmeasurements at low pH—It is directly determined from the absorbanceratios shown in FIG. 6 ((e₃/e₂)=(₂₅₀A/₂₃₄A) ⁻¹)); (c) using (e₃/e₂)values determined at low pH, Eq. (20) can be used with paired [CO₃²⁻]_(T) and R observations to directly determine (_(CO3)β₁)/(e₂) and e₁.The results of such analyses, using each of the data sets that wereemployed to develop Eq. (11) and Eqs. (17) through (19), are given inTable 5 and are depicted graphically in FIGS. 7 a and 7 b. FIG. 8 showsthe residuals, ([CO₃ ²⁻]_(T))_(observed)−([CO₃ ²]hd TT/_(predicted)),for each least squares analysis using Eq. (20). It should be noted thatthe residuals shown in FIG. 8 are derived from three independentsources. One of these is the absorbance ratios that are used to predict[CO₃ ²⁻]_(T) via Eq. (20), and the others are the alkalinity and pHmeasurements that are used to derive “observed” values of [CO₃ ²⁻]_(T).Of these three types of measurements, there is reason to suspect thatpotentiometric pH measurements may constitute the greatest source ofscatter seen in FIG. 8. Since field measurements of [CO₃ ²⁻]_(T) basedon absorbance ratios are independent of alkalinity and pH, the precisionof [CO₃ ²⁻]_(T) measurements obtained via absorbance spectroscopy shouldbe considerably better than that which is depicted in FIG. 8. Theresults shown in FIG. 8 suggest that Eq. (20) can be used tosatisfactorily predict [CO₃ ²⁻]_(T) over a wide range of conditions inseawater. Preliminary assessment of the relative contributions ofpotentiometric and spectrophotometric contributions to the residuals inFIG. 8 suggest that the relative standard deviation for Eq. (20)calculations of [CO₃ ²⁻]_(T) is on the order of 2% or less. The bestleast squares descriptions for the salinity dependencies of theparameters in Eq. (20) for 20≦S≦36 are given as follows:log {(_(CO3)β₁)/(e ₂)}=6.087−8.495×10⁻² S+9.360×10⁻⁴  (21)e ₁=0.2215−5.554×10⁻⁴ S+8.440×10⁻⁵  (22)(e ₃ /e ₂)=3.061−8.730×10⁻² S+9.363×10⁻⁴  (23)

TABLE 4 e₁ (₂₅₀ε_(PbCO3)/₂₃₄ε_(PbCO3)), e₂ (₂₅₀ε_(Pb)/₂₃₄ε_(PbCO3)) ande₃ (₂₃₄ε_(Pb)/₂₃₄ε_(PbCO3)) as a function of salinity at 25° C. Salinitye₁ e₂ e₃ 35.87 0.2938 0.8009 0.9068 34.50 0.2890 0.7841 0.9131 32.500.2828 0.7619 0.9263 30.00 0.2766 0.7381 0.9498 27.50 0.2721 0.71820.9825 25.00 0.2694 0.7024 1.0262 22.50 0.2687 0.6908 1.0835 20.000.2701 0.6838 1.1588

TABLE 5 Best fit log{(_(CO3)β₁)/(e₂)} and e₁ results obtained using Eq.(20). The (e₃/e₂) values used in Eq. (20) are the reciprocals of the₂₅₀A/₂₃₄A values given in Table 3. Salinity log{(_(CO3)β₁)/(e₂)} e₁35.87 4.202 0.2957 34.50 4.304 0.3217 32.50 4.328 0.2909 30.00 4.4260.2867 27.50 4.400 0.2582 25.00 4.542 0.2609 22.50 4.650 0.2538 20.004.773 0.2449

It should be emphasized that the equation (22) and (23)characterizations of e₁ and e₃/e₂ were obtained through fits involvingequation (20) and did not involve the e₁, e₂ and e₃ characterizationsgiven in equations (17) through (19).

Using equations (20) through (23) carbonate ion concentrations can bemeasured via procedures that closely follow those employed by Claytonand Byrne (1993) for measurements of seawater pH:

-   -   Seawater obtained from Niskin bottles is transferred directly to        10 cm quartz cuvettes without exposure to the atmosphere        (samples are not filtered).    -   Samples in the quartz cuvettes (total volume ˜30 cm³) are        thermostatted at 25° C.    -   The thermostatted sample is directly used as a reference        (baseline) solution (_(λ)A=0 at all wavelengths). Note, in this        case, that absorbance contributions from UV active species such        as nitrate are identical in the reference (baseline) measurement        and in subsequent measurements wherein Pb(II) is added to the        solution. As such, the absorbance contributions of such species        (nitrate, organics etc.) are eliminated via baseline        subtraction.    -   A 1 mM stock solution of PbCl₂ is added to the cuvette (˜0.22        cm³ addition) whereupon the Pb(II) concentration is        approximately 7.5 μM.    -   Absorbances are measured at 234, 250 and 350 nm, and the        absorbance at λ=350 nm is used to correct for any small baseline        changes induced by manipulation of the cuvette.    -   Using the absorbance ratio obtained through this protocol        (₂₅₀A/₂₃₄A), and the salinity dependent coefficients in        equations 21-23, equation 20 is used to calculate log [CO₃        ²⁻]_(T).

It should be noted that the relatively high lead concentrations used formeasurements of carbonate ion concentrations preclude significantcomplexation by dissolved and particulate organics in seawater. Sinceionic Pb(II) in seawater is not significantly volatile, the principalsafety concern surrounding the manipulation of solutions enriched inlead is avoidance of inadvertent ingestion. This should especially beborne in mind with respect to handling of the 1 mM stock Pb titrantsolution and safe disposal of waste solutions.

Determinations of Seawater Salinity

Salinity measurements required for calculations of log{(_(CO3)β₁)/(e₂)}, e₁, and (e₃/e₂) are commonly available from eithershipboard or in situ conductimetric observations. When such is not thecase, however, salinity can be calculated from absorbance ratios usingthe following relationship that is based on the data shown in FIG. 5 foracidified seawater:S=−8.76+45.15R+6.092R ²  (24)

where R=₂₅₀A/₂₃₄A and 20≦S≦36. The standard deviation for Eq. (24)estimates of salinity is ±0.06 salinity units. Thus, Eq. (24) providesseawater salinity estimates that are precise to approximately 0.2% overthe normal salinity range of seawater.

Thus, salinity can be measured from absorbance ratio observations afterPb(II) is added to an acidic seawater solution. Other metals can be usedto obtain sensitive salinity measurements over various ranges ofsalinity. The work reported herein was performed at 25° C. It followsthat the salinity measurements could be performed at other temperaturesafter suitable laboratory calibrations. This would allow salinitymeasurements and/or carbonate measurements to be obtained withoutthermostatting. Such calibrations facilitate carbonate measurements insitu over a range of temperatures.

Potentiometric pH Measurements and Systems Calibration

The potentiometer measures the millivolts developed by a pH electrode vsa reference electrode. The solution measured consists of the aqueousmedium plus a small amount of indicator. The pH of this mixture isassessed by both the spectrophotometer and the potentiometer. Thepotentiometric pH is determined by an equation of the following form:pH=a+b·(millivolts)  (25)

The “b” term in the above equation (Eq. 25) is an electrode slope thatdepends on temperature. The “a” term depends on a number of variablesand can be determined via calibrations. In the present instance, thepotentiometric system (or the calibration system) is “told” the pH of agiven solution (i.e., the pH value is determined spectrophotometrically)whereupon the constant “a” is determined from the above equation. Insuch a system it is also possible to have an additional pump so thataqueous solution, indicator and acid are combined. In this case pH isagain measured spectrotometrically and the above equation can be used tosolve for both slope “b” and intercept “a”.

CONCLUSIONS

In the absence of direct spectrophotometric determinations of [CO₃²⁻]_(T) as described above, [CO₃ ²⁻]_(T) must be calculated frommeasurements of either total dissolved inorganic carbon or totalalkalinity combined with either pH or CO₂ fugacity. Whilespectrophotometric pH measurements are rapid, with acquisition rates onthe order of seconds, measurements of dissolved inorganic carbon, totalalkalinity and CO₂ fugacity generally require several minutes (as alower bound). Thus, the spectrophotometric procedures for measurementsof carbonate ion concentrations described in this work, and those forspectrophotometric pH analysis (Liu, et al., 2006), are unique in theirsuitability for prompt in situ analysis. Although, in situ analysis willrequire evaluations of the influence of temperature on the variousparameters in Eq. (20), the work of Soli et al. (2008) showed that theinfluence of temperature on log(_(CO3)β₁) is quite small. Since theinfluence of temperature on molar absorptivity ratios should berelatively minor, it is likely that Eq. (20) can easily be extended toinclude analysis at in situ conditions. With respect to both in situ andlaboratory analysis, spectrophotometric pH and [CO₃ ²⁻]_(T) measurementprocedures can be distinguished from those required for C_(T) withrespect to both procedural and instrumental simplicity. In contrast tothe equipment required for state of the art C_(T) analyses,spectrophotometers are standard equipment in a wide variety of researchand teaching laboratories.

As has been the case for spectrophotometric measurements of pH, itshould be anticipated that the parameters required for quantitative [CO₃²⁻]_(T) measurements (e.g., Eqs. (21) through (24)) will periodically bereevaluated and refined. This process will include, in particular,comparisons obtained through ship-based oceanic carbon systemexpeditions wherein the thermodynamic consistency of all measurable CO₂system parameters is commonly evaluated (Clayton et al., 1995; Lee etal., 2000). It should be emphasized in this case that any futurerevisions in characterizations of log {(_(CO3)β₁)/(e₂)}, e₁ and (e₃/e₂)will allow refinements, with improved accuracy, of archived [CO₃ ²⁻]_(T)data. As long as data are recorded as R—S pairs (i.e., absorbance ratiosand salinity), all calculations of [CO₃ ²⁻]_(T) are amenable toquantitative reassessment. As such, observations of Pb(II) absorbanceratios provide a molecularly-based index of carbonate ion concentrationsin seawater.

The procedures described in this work are suitable for rapid,quantitative assessments of calcite and aragonite saturation states inseawater. Since the solubility products of calcite and aragonite in S=35seawater are approximately 10^(−6.367) and 10^(−6.186) (Millero, 2007),and the total calcium concentration is 0.0103 mol/kg at salinity 35, thecarbonate ion concentrations for saturation with calcite and aragoniteare 41.7 μmol/kg and 63.3 μmol/kg. The log _(CO3)β₁ results given by Eq.(11) (log _(CO3)β₁=4.106 at S=35) show that inorganic Pb(II) ispartitioned equally between PbCO₃ ⁰ and lead chloride complexes when[CO₃ ²⁻]_(T)=78.3 mmol/kg. Thus, the procedures described in this workare well suited to measurement of CaCO₃ saturation states both below andwell above the saturation levels of calcite and aragonite.

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It will be seen that the advantages set forth above, and those madeapparent from the foregoing description, are efficiently attained andsince certain changes may be made in the above construction withoutdeparting from the scope of the invention, it is intended that allmatters contained in the foregoing description or shown in theaccompanying drawings shall be interpreted as illustrative and not in alimiting sense.

It is also to be understood that the following claims are intended tocover all of the generic and specific features of the invention hereindescribed, and all statements of the scope of the invention which, as amatter of language, might be said to fall therebetween. Now that theinvention has been described,

1. A method of performing in situ continuous autonomous potentiometricpH measurements for extended periods of time comprising the steps of:recording a plurality of potentiometric pH measurements using apotentiometric pH measurement device; autonomously activating aspectrophotometric calibration system positioned within the devicewherein the system is activated by an event trigger or a defined timeinterval wherein the event trigger is a sudden pH change outside thetypical deviations of pH recordings; calibrating the potentiometric pHmeasurement device comprising the steps of: obtaining spectrophotometricpH measurements from a spectrophotometer located within the devicecomprising the steps of: sampling an aqueous medium; collecting datarepresentative of temperature and salt concentration of the sampledaqueous medium; combining a sulfonephthalein indicator and a firstportion of the sampled aqueous medium; delivering the combined medium toan optical cell; obtaining an absorbance spectrum from the combinedmedium; calculating solution pH of the sampled aqueous medium utilizingthe optical absorbance ratios at multiple wavelengths, the temperature,and the salt concentration; obtaining a potentiometric pH measurement ofa second portion of the aqueous medium substantially contemporaneouslywith the spectrophotometric pH measurements; calculating the deviationbetween the potentiometric pH measurement and the spectrophotometric pHmeasurement; and adjusting the potentiometric pH measurement deviceusing the calculated deviation, wherein the adjusting calibrates thepotentiometric pH measurement device; and recording the calibratedpotentiometric pH measurements for an extended period of time until anevent trigger or preset time interval autonomously activates thespectrophotometric calibration system to calibrate the device.
 2. Themethod of claim 1 wherein measurement frequency of the potentiometric pHmeasurement is 60 Hz or below.